1. Chapter 20 Electrochemistry
20.1 Introduction to Electrochemistry

2. Electrochemistry
The branch of chemistry that deals with electricity-related applications of oxidation-reduction reactions.
Electrochemical Cells:
A system of electrodes and electrolytes in which either chemical reactions produce energy or an electrical current produces chemical change

3. Components of Electrochemical Cells
Cu Electrode
Cathode- where reduction takes place
Conducting Wire
Electrolyte Sol’n ZnSO4
Electrolyte Sol’n CuSO4
Half-Cell: a single electrode immersed in a solution of its ions
Electrode: conductor used to establish electrical contact with a nonmetallic part of the circuit.
Zn Electrode
Anode- where oxidation takes place

4. Half-Cell: a single electrode immersed in a solution of its ions
Cu Electrode
Cathode- written as
Cu+2/Cu
Half-Cell: a single electrode immersed in a solution of its ions
Overall Cell Written as:
anode | cathode
Zn | Cu
Zn Electrode
Anode- written as
Zn+2/Zn

6. Chapter 20 Electrochemistry
20.2 Voltaic Cells

7. Voltaic / Galvanic Cell Rxns that produce voltage spontaneously
Electrochemistry
Porous barrier which prevents the spontaneous mixing of the aqueous solutions in each compartment, but allows the movement of ions in both directions to maintain electrical neutrality
Voltaic / Galvanic Cell
Rxns that produce voltage spontaneously
A chemical rxn that results in a voltage due to a transfer of electrons

8. Batteries Two or more dry voltaic cells Zinc-Carbon Battery
Zn → Zn+2 + 2e-
2MnO2 + H2O + 2e- → Mn2O3 + 2OH -

9. Batteries Alkaline Battery- no carbon rod, smaller
Zn + 2OH - → Zn(OH)2 + 2e-
2MnO2 + H2O + 2e- → Mn2O3 + 2OH-

10. Batteries Mercury Battery- no carbon rod, smallest
Zn + 2OH - → Zn(OH)2 + 2e-
HgO + H2O + 2e- → Hg + 2OH -

11. Rxns that turn chemical energy into electrical energy
Fuel Cells
Cathode: O2 + 2H2O + 4e- → 4OH –
Anode: 2H2 + 4OH – → 4e- + 4H2O
Net: 2H2 + O2 → 2H2O
A voltaic cell where reactants are constantly supplied and products are removed.
Rxns that turn chemical energy into electrical energy

13. Corrosion Formation of Rust: 4Fe (s) + 3O2 (g) + xH2O → 2Fe2O3∙xH2O
Anode: Fe (s) → Fe+2 (aq) + 2e-
Cathode: O2 (g) + 2H2O (l) + 4e- → 4OH –

14. Prevention of Corrosion
Galvanizing
Process by which iron or any metal is coated with zinc. Cathodic Protection
Since zinc is more easily oxidized, it is a sacrificial anode.

15. Electrode Potentials
Reduction Potential: the tendency for the half-reaction to occur as a reduction half-reaction in an electrochemical cell.
Electrode Potential: the difference in potential between an electrode and its solution
Potential Difference (Voltage): a measure of the energy required to move a certain electric charge between the electrodes, measured in volts.
Standard Electrode Potential (E°): a half-cell measured relative to a potential of zero for the standard hydrogen electrode (SHE)

16. Standard Electrode Potential, E°
Positive E° means hydrogen is more willing to give up its electron, so positive reduction potentials are favored. Naturally occurring rxns have a positive value.
E° cell = E° cathode - E° anode
Negative E° means the metal electrode is more willing to give up its electron, this is not favored. These rxns prefer oxidation over reduction.

17. Standard Electrode Potential, E°
When a half-cell is multiplied by a constant (for balancing) the E° value is NOT multiplied!
When a rxn is reversed (flipped) the sign of the E° value switches.
In a voltaic cell, the half-rxn with the more negative standard electrode potential is the anode, where oxidation occurs.

18. Because this is a spontaneous process:
Cell Potential
The potential voltage a rxn can produce.
Reduction potentials
Cu e-  Cu Eo = .34 V
Ag+ + e-  Ag Eo = .80V
Since both rxns are reduction, one must be oxidation, flip it, positive voltage must result from spontaneous rxns
Because this is a spontaneous process:
(Ag+ + e-  Ag) x Eo = .80V
Cu  Cu e Eo = -.34 V
Cu + 2Ag+ Cu Ag Eo = .46 V

19. Because this is nonspontaneous process:
Cell Potential
The potential voltage a rxn can produce.
Na+ + e-  Na Eo = V
Nonspontaneous, must end in negative voltage. Flip one to become oxidation.
** Fuel Cell!
Cl2 + 2e-  2Cl- Eo = 1.36 V
Because this is nonspontaneous process:
(Na+ + e-  Na) x 2 Eo = V
2Cl-  Cl e- Eo = V
2Na Cl-  2Na + Cl2 Eo = V

20. Chapter 20 Electrochemistry
20.3 Electrolytic Cells

21. Rxns that require an energy source to react
Electrochemistry
When electric voltage is used to produce a redox reaction, it is called electrolysis
Electrolytic Cell
Rxns that require an energy source to react

22. Batteries
Car Battery- rechargeable b/c the alternator reverses the ½ rxns and regenerates the reactants.
Discharge Cycle Rxn:
Pb + PbO2 + 2H2SO4 → 2PbSO4 + 2H2O

23. Electroplating
An electrolytic process in which a metal ion is reduced and a solid metal is deposited on a surface
Typically, an inactive metal is able to be ionized and then deposited on the surface of a more active metal to prevent corrosion.
Anode
Silver ions are reduced at the cathode:
Ag+ + 1e- → Ag
Silver atoms are oxidized at the anode:
Ag → Ag + + 1e-
Cathode

24. Voltaic vs. Electrolytic
If the positive battery terminal is attached to the cathode of a voltaic cell, and the negative terminal is attached to the anode, the flow of electrons will change directions.
Electrolytic cells need the electrodes attached to a battery, where voltaic is its own source of electrical power.
Voltaic = spontaneous
chemical energy → electrical energy
Electrolytic = non-spontaneous
electrical energy → chemical energy

25. Electrolysis
Anode: 6H2O → O2 + 4e- + 4H3O+
Cathode: 4H2O + 4e- → 2H2 + 4OH –
Using a current to generate a redox reaction which otherwise would have a negative cell potential. i.e. electroplating & rechargeable batteries.